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- Dec 23, 2009
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Hi all,
My current understanding of van der waals and volume/pressure:
-gas molecules occupy no volume under ideal conditions.
-the van der waals equation is correcting for devations due to real gases.
-if there are attractive forces, pressure of a REAL gas decreases (since there are less collisions with the containter wall) so to correct you add the a(n2/V2) term to real pressure to find pressure under ideal conditions
My problem comes with volume. Under ideal conditions, V of the PV=nRT equation is the total volume of the container (because ideally the gas molecules weigh nothing). In reality gas molecules do weigh some amount so we correct by subtracting the nb term. The goal of the ideal gas law is to show us ideal conditions. Why then do we even correct for the volume of molecules at all? If you subract the nb term you get a volume of the container less than ideal conditions that seems like REAL conditions right? The pressure correction makes complete sense because you want to correct it back to ideal conditions so you add the attraction value to make the pressure in the equation what it would be under ideal conditions. But if you subtract the volume of the molecules, wouldn't that term (V-nb) be what you should get under real conditions? I guess my question is why is that the proper way to correct for volume deviations from ideality (isn't ideality just V itself?)?
I hope my questions make sense, thank you for any help!
My current understanding of van der waals and volume/pressure:
-gas molecules occupy no volume under ideal conditions.
-the van der waals equation is correcting for devations due to real gases.
-if there are attractive forces, pressure of a REAL gas decreases (since there are less collisions with the containter wall) so to correct you add the a(n2/V2) term to real pressure to find pressure under ideal conditions
My problem comes with volume. Under ideal conditions, V of the PV=nRT equation is the total volume of the container (because ideally the gas molecules weigh nothing). In reality gas molecules do weigh some amount so we correct by subtracting the nb term. The goal of the ideal gas law is to show us ideal conditions. Why then do we even correct for the volume of molecules at all? If you subract the nb term you get a volume of the container less than ideal conditions that seems like REAL conditions right? The pressure correction makes complete sense because you want to correct it back to ideal conditions so you add the attraction value to make the pressure in the equation what it would be under ideal conditions. But if you subtract the volume of the molecules, wouldn't that term (V-nb) be what you should get under real conditions? I guess my question is why is that the proper way to correct for volume deviations from ideality (isn't ideality just V itself?)?
I hope my questions make sense, thank you for any help!
